Figure a ∇plots the first ionization energies of the...

Figure a $\nabla$ plots the first ionization energies of the elements in period $3 .$ Figure $\mathrm{b}$ > lists some of the binary compounds that form between fluorine and the elements in period $3 .$ Figure $\mathrm{c}$ > plots the electron affinities of the elements in period 3 . Use the information provided in the figures to answer the following questions: a. In general, first ionization energy increases as we move from left to right across the third period. Explain. b. There is a decrease in first ionization from $\mathrm{Mg}$ to $\mathrm{Al}$. There is an increase in electron affinity from $\mathrm{Mg}$ to $\mathrm{Al}$. Explain. c. The first ionization energy of $\mathrm{S}$ is slightly less than that of $\mathrm{P}$. Explain. d. After reviewing the information in Figures $\mathrm{a}, \mathrm{b},$ and $\mathrm{c},$ Student X offers the following hypothesis: The number of fluorine atoms that bond to a metal atom is equal to the number of valence electrons of the metal. The number of fluorine atoms that bond to a nonmetal atom is equal to 8 minus the number of valence electrons of the nonmetal. Student Y offers the following hypothesis: In general, there is an inverse relationship between the electron affinities of the period 3 elements and the number of fluorine atoms that bond to the metal or nonmetal atoms. Do you agree with Student X's hypothesis, with Student Y's hypothesis, both of them, or neither of them? Explain. e. Write the formula for a compound that forms between silicon and flourine. Does this compound support Student X's hypothesis, Student Y's hypothesis, or neither hypothesis? Explain.

Figure a $\nabla$ plots the first ionization energies of the elements in period $3 .$ Figure $\mathrm{b}$ > lists some of the binary compounds that form between fluorine and the elements in period $3 .$ Figure $\mathrm{c}$ > plots the electron affinities of the elements in period 3 .
Use the information provided in the figures to answer the following questions:
a. In general, first ionization energy increases as we move from left to right across the third period. Explain.
b. There is a decrease in first ionization from $\mathrm{Mg}$ to $\mathrm{Al}$. There is an increase in electron affinity from $\mathrm{Mg}$ to $\mathrm{Al}$. Explain.
c. The first ionization energy of $\mathrm{S}$ is slightly less than that of $\mathrm{P}$. Explain.
d. After reviewing the information in Figures $\mathrm{a}, \mathrm{b},$ and $\mathrm{c},$ Student X offers the following hypothesis: The number of fluorine atoms that bond to a metal atom is equal to the number of valence electrons of the metal. The number of fluorine atoms that bond to a nonmetal atom is equal to 8 minus the number of valence electrons of the nonmetal.
Student Y offers the following hypothesis: In general, there is an inverse relationship between the electron affinities of the period 3 elements and the number of fluorine atoms that bond to the metal or nonmetal atoms.
Do you agree with Student X's hypothesis, with Student Y's hypothesis, both of them, or neither of them? Explain.
e. Write the formula for a compound that forms between silicon and flourine. Does this compound support Student X's hypothesis, Student Y's hypothesis, or neither hypothesis? Explain.

Key Concepts

Ionization Energy

Ionization energy is the amount of energy required to remove the most loosely held electron from an isolated gaseous atom. It generally increases across a period because of the increasing effective nuclear charge which draws the electrons closer, making them harder to remove. Variations in ionization energy between adjacent elements often arise from differences in electron configurations and subshell stability, particularly when moving from a filled or half-filled subshell to one that is not optimally filled.

Periodic Trends

Periodic trends refer to the systematic variations in elements' properties as one moves across a period or down a group in the periodic table. These trends—such as changes in ionization energy, atomic radius, electronegativity, and electron affinity—occur due to variations in nuclear charge and electron shielding. An understanding of these trends helps explain why certain properties generally increase or decrease across periods and down groups, despite occasional irregularities due to electron configuration effects.

Electron Affinity

Electron affinity is the energy change that occurs when an electron is added to a neutral atom in the gas phase, and it is a measure of an atom's tendency to accept an electron. In general, electron affinity becomes more exothermic across a period as atoms become more eager to fill their valence shell for achieving a stable electronic configuration. However, anomalies can occur due to electron-electron repulsions in more compact electron arrangements.

Subshell Electron Configuration and Stability

Electron configurations determine the distribution of electrons in an atom’s orbitals, which in turn affects its chemical properties. Stable configurations, such as those with half-filled or fully filled subshells, often require extra energy to disrupt. These stability effects explain deviations in periodic trends such as anomalously low ionization energies or unexpectedly high electron affinities between elements with similar valence electron counts.

Chemical Bonding and Compound Formation

Chemical bonding involves the interaction between atoms so that they achieve a stable electronic configuration, often modeled by the octet rule. In forming compounds, metals tend to lose electrons and nonmetals tend to gain electrons, and the stoichiometry of these compounds can often be predicted by the number of valence electrons an atom has. Understanding the electron arrangements helps in predicting the formulas of binary compounds, as well as in evaluating hypotheses related to bonding patterns and the ratios of atoms in compounds.

Transcript

00:02 Using the information presented to us in figures a, b, and c, let's answer the following questions.

00:10 For a, the first ionization energy increases as we move from left to right across the third period.

00:18 So ionization energy increases moving across a period as the nuclear.

00:38 Charge increases moving from left to right across a period, creating a greater attraction for electrons, therefore increasing the ionization energy.

01:27 For b, we're asked to look at magnesium and aluminum.

01:32 So let's first look at, our electron configurations for magnesium and aluminum.

01:39 Magnesium is ending in 3s1 or sorry 3s2 and aluminum ending in 3p1.

01:55 We're first asked to explain the ionization energy difference.

02:03 So the magnesium orbitals are completely filled.

02:26 So more energy will be required to remove an electron.

02:43 So therefore, the ionization energy of magnesium is higher than aluminum.

03:02 Now let's look at this in terms of electron affinity.

03:12 Sorry, this would be 3p1 here.

03:23 So magnesium is stable.

03:30 With its complete orbits.

03:41 So its electron affinity is near zero, whereas aluminum has an incomplete 3p orbital.